Energy Levels

In all atoms other then hydrogen, electrons in differentorbitals have varying energy levels with those furthest away from the nucleus having the greatest ones. An electron’s energy level is not only determines by its distance from the nucleus but all by the presence of other electrons. Just as the positive nucleus attracts it, the negative charges of its neighbouring electrons repel it. This phenomenon does not occur in Hydrogen, since it has a single electron but the repulsion of outer electrons by inner ones occurs in all other atoms. Part of the attractive force of the nucleus is disrupted by the repelling force of the inner electrons and we say they shield the outermost electrons from the nucleus. As a result, the electrons are held less tightly which has a profound effect on the trends visible throughout the periodic table.The closer an electron is to the nucleus, the lower its energy level and the orbitals fill such that the next added electron will occupy the lowest energy levelThe energy levels of different orbitals are indicated. Though the energy levels in orbitals of Hydrogen depend only on the value of n, in other atoms the energy depends on both n and l. Note the crossover ofenergies from 1 orbital to another, particularly the fact that 4s has a lower energy level than 3d.

Electron Configuration

The electron configuration for an element is the arrangement of electrons in the orbits (or shells) of a neutral atom. Shells closer to the nucleus have higher binding energy. The electron configuration is described by a notation that lists the subshell symbols, one after another. Each symbol has a subscript on the right indicating the number of electrons occupied in that subshell. For example, the electron configuration of Oxygen (atomic number 8) is 1s^22s^22p^4. The first number seen in the electron configuration for oxygen (1) represents the main energy level. There are 7 main energy levels, same as the 7 rows of the periodic table. Each level contains a number of sublevels which contain a number of orbitals. Each orbital can contain 2 electrons each. Sublevel s contains 1 orbital. Sublevel p contains 3 orbitals. Sublevel d contains 5 orbtials and sublevel f contains 7 orbitals. Knowing these rules, finding the electron configuration of an element will be simple.

Can you figure out the electron configuration of element K (Potassium)? Since Potassium contains 19 electrons, firstly, two electrons will fill the 1s orbital, leaving 17 electrons. Secondly, the next two electrons will fill the 2s orbital, leaving 15 electrons. The 2p orbital is the next available energy level and can hold six electrons. This leaves 9 electrons to be filled. The next two electrons will fill the 3s orbital, leaving 7 electrons. The 3p is the next energy level and like the 2p orbital, it can also hold six electrons, leaving 1 electron left. This 1 electron will fill one-half of the 4s orbital. Therefore, the electron configuration of Potassium is 1s^2 2s^2 2p^6 3s^2 3p^6 4s^1.

Electron Configuration Shortcut(Noble Gas Simplification)At times, there will be elements where the electron configuration will be very long. However, there is a shortcut of writing electron configuration! As mentioned before, writing the electron configuration of an element involves determining all the orbitals that are filled until every electron has filled an orbital. We learnt that oxygen's electron configuration was:1s^22s^22p^4. It can also be written as: [He] 2s^2 2p^4. This is another way for writing electron configuration. The steps in determining this method are:1. Firstly, start at the first noble gas before the element you are going to name. In Oxygen's case, Helium is the first noble gas before it. You would write it as [He]. 2. Then you would follow the rest of the rules of determining its electron configuration. Since [He] acts as a replacement for 1s^2, everything else stays the same.

Can you figure out the electron configuration of Potassium using this method? Firstly, look at the noble gas before Potassium. Conveniently, it's found one element behind Potassium, Argon. This would be written as [Ar].Secondly, determine how many electrons are left. Since only 1 electron is now left over, we know that it has to fill one-half of the 4s orbital due to the fact that Argon occupies the previous orbitals. Therefore, using this method, another way of writing Potassium's electron configuration is [Ar] 4s^1.

The Aufbau Principle

The Aufbau Principle was a key component to Niels Bohr's original concept of electron configuration. It can be stated that a maximum of two electrons are put into orbitals in the order of increasing orbital energy: the lowest-energy orbitals are filled before electrons are placed in higher-energy orbitals.

Orbitals are filled in the order of increasing n+l;

Where two orbitals have the same value of n+l, they are filled in order of increasing n.

No two electrons in the atom will share the same four quantum numbers n, l, m, and s. This means that the two electrons must have opposite (paired spins).

Pauli Exclusion Principle

The Pauli Exclusion Principle is part of the Aufbau Principle. The Pauli Exclusion Principle is a quantum mechanical principle which states that if two electrons are in an orbital, they must have opposite spins. As a result, no two electrons in the atom will share the same set of four quantum numbers.

This diagram shows the correct orientation This diagram is incorrect. Electrons must spin in of electrons occupying the same orbital. opposite directions.

Hund’s Rule

Hund's Rule is also a part of the Aufbau Principle. Hund's Rule states that if two or more orbitals of equal energy are available, electrons will occupy them singly before filling them in pairs. Electrons will occupy individual orbitals before they pair up.

Applications

Knowledge of electron configuration has a number of applications, one of the most important being its ability to accurately explain periodic trends that manifest in Mendeleev's table. One such trend is atomic radius which increases down a group and decreases across a period. Hence Sodium will have a larger atomic radius than Lithium and this can be easily explained by considering the electron configuration. Lithium has its outermost electron in a 2s orbital while Sodium’s outermost electron occupies 3s. Similarly Potassium’s electrons are in 4s and Rubidium’s in 5s. As you go down a group the shells electrons occupy become larger and so the electrons are further from the nucleus, resulting in a greater atomic radius. The trend across the period is slightly trickier to explain but not beyond our abilities. As noted previously, inner electrons shield outer ones from the some of the nucleus’s attractive forces. As you move across a period however, electrons are added to the outer shells, not the inner ones, and the number of protons in the nucleus continues to increase. Electrons in the same shell are roughly the same distance from the nucleus as its neighbours and are ineffective at shielding them from the effects of the nucleus. Hence the attractive force of the nucleus on the outer electrons increases across a period, drawing the atoms closer to the atom’s centre and decreasing the atomic radius.